Why Study Electrochemistry? • Batteries • Corrosion • Industrial production of chemicals such as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group Review of Terminology for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number. • REDUCTION—gain of electron(s); decrease in oxidation number. • OXIDIZING AGENT—electron acceptor; species is reduced. • REDUCING AGENT—electron donor; species is oxidized. Electrochemical Cells • An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. • Indirect redox. • There are two types: one product favored, the other reactant favored. Batteries are voltaic cells Electrochemical Cells • Product favored reaction -> voltaic or galvanic cell -> electric current produced -> e.g., batteries • Reactant favored reaction -> electrolytic cell -> electric current used to cause chemical change. -> e.g., used to obtain pure metals Basic Concepts of Electrochemical Cells Anode Cathode CHEMICAL CHANGE ---> ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” •Zn is oxidized and is the reducing agent Zn(s) ---> Zn2+(aq) + 2e•Cu2+ is reduced and is the oxidizing agent Cu2+(aq) + 2e- ---> Cu(s) CHEMICAL CHANGE ---> ELECTRIC CURRENT Electrons are transferred from Zn to Cu2+, but there is no useful electric current. Oxidation: Zn(s) ---> Zn2+(aq) + 2eReduction: Cu2+(aq) + 2e- ---> Cu(s) -------------------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s) CHEMICAL CHANGE ---> ELECTRIC CURRENT •To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery. Zn --> Zn2+ + 2e- Cu2+ + 2e- --> Cu Oxidation Anode Negative Reduction Cathode Positive <--Anions Cations--> •Electrons travel thru external wire. •Salt bridge allows anions and cations to move between electrode compartments. Terms Used for Voltaic Cells Figure 20.3 CELL POTENTIAL, E 1.10 V Zn and Zn2+, anode 1.0 M 1.0 M Cu and Cu2+, cathode • Electrons are “driven” from anode to cathode by an electromotive force or emf. • For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. CELL POTENTIAL, E • For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. • This is the STANDARD CELL POTENTIAL, Eo • —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C. Calculating Cell Voltage • Balanced half-reactions can be added together to get overall, balanced equation. Zn(s) ---> Zn2+(aq) + 2eCu2+(aq) + 2e- ---> Cu(s) -------------------------------------------Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s) If we know Eo for each half-reaction, we could get Eo for net reaction. CELL POTENTIALS, o E Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL, SHE. 2 H+(aq, 1 M) + 2e- <----> H2(g, 1 atm) o E = 0.0 V Zn/Zn2+ half-cell hooked to a SHE. Eo for the cell = +0.76 V Negative electrode Supplier of electrons Zn --> Zn2+ + 2eOxidation Anode Positive electrode Acceptor of electrons 2 H+ + 2e- --> H2 Reduction Cathode Reduction of H+ by Zn Figure 20.10 Overall reaction is reduction of H+ by Zn metal. Zn(s) + 2 H+ (aq) --> Zn2+ + H2(g) Eo = +0.76 V Therefore, Eo for Zn ---> Zn2+ (aq) + 2e- is +0.76 V Zn is a (better) (poorer) reducing agent than H2. Cu/Cu2+ and H2/H+ Cell Eo = +0.34 V Positive Acceptor of electrons Cu2+ + 2e- --> Cu Reduction Cathode Negative Supplier of electrons H2 --> 2 H+ + 2eOxidation Anode 2+ Cu/Cu and H2 + /H Cell Overall reaction is reduction of Cu2+ by H2 gas. Cu2+ (aq) + H2(g) ---> Cu(s) + 2 H+(aq) Measured Eo = +0.34 V Therefore, Eo for Cu2+ + 2e- ---> Cu is +0.34 V Zn/Cu Electrochemical Cell + Anode, negative, source of electrons Cathode, positive, sink for electrons Zn(s) ---> Zn2+(aq) + 2eEo = +0.76 V Cu2+(aq) + 2e- ---> Cu(s) Eo = +0.34 V --------------------------------------------------------------Cu2+(aq)+Zn(s) ---> Zn2+(aq)+Cu(s) Eo (calc’d) = +1.10 V Uses of Eo Values • Organize halfreactions by relative ability to act as oxidizing agents • Table • Use this to predict cell potentials and direction of redox reactions. TABLE OF STANDARD REDUCTION POTENTIALS oxidizing ability of ion Eo (V) Cu2+ + 2e- Cu +0.34 2 H+ + 2e- H2 0.00 Zn2+ + 2e- Zn -0.76 reducing ability of element Use Reduction Potentials to decide on cell. Cu|Cu2+ and Ag|Ag+ The Cu|Cu2+ and Ag|Ag+ Cell More About Calculating Cell Voltage Assume I- ion can reduce water. 2 H2O + 2e- ---> H2 + 2 OHE˚= 2 I- ---> I2 + 2eE˚= ---------------------------------------------------------------------2 I- + 2 H2O --> I2 + 2 OH- + H2 E˚= Assuming reaction occurs as written, Minus E˚ means rxn. occurs is _____________. Electrolysis Using electrical energy to produce chemical change. Sn2+(aq) + 2 Cl-(aq) ---> Sn(s) + Cl2(g) SnCl2(aq) Cl2 Sn Electrolysis of Aqueous NaOH Electric Energy ----> Chemical Change Anode (+) 4 OH- ---> O2(g) + 2 H2O + 4eCathode (-) 4 H2O + 4e- ---> 2 H2 + 4 OHEo for cell = _______ Electrolysis Electric Energy ---> Chemical Change • Electrolysis of molten NaCl. • Here a battery + “pumps” electrons Anode from Cl- to Na+. • Polarity of electrodes is reversed from batteries. electrons BATTERY Cathode Cl- Na+ Electrolysis of Molten NaCl Figure 20.14 Electrolysis of Molten NaCl electrons Anode (+) BATTERY 2 Cl- ---> Cl2(g) + 2e- Eo = + Anode Cathode Cl- Na+ Cathode (-) Na+ + e- ---> Na Eo = Eo for cell (in water) = ___________ External energy needed because Eo is (-). Electrolysis of Aqueous NaCl Anode (+) 2 Cl- ---> Cl2(g) + 2eCathode (-) 2 H2O +2e- -->H2 + 2 OHEo for cell = -2.19 V Note that H2O is more easily reduced than Na+. Also, Cl- is oxidized in preference to H2O because of kinetics. Electrolysis of Aqueous NaCl Cells like these are the source of NaOH and Cl2. In 1995: 25.1 x 109 lb Cl2 and 26.1 x 109 lb NaOH Also the source of NaOCl for use in bleach. Is E˚ related to ∆G? YES! But you don’t have to worry about the equation….just remember +E° is spontaneous = product favored= Electrochemical Cell (Galvanic, Voltaic) -E° is not spontaneous = reactant favored = Electrolytic Cell